2063
Solvent Effects on the Reactions of Stabilized Carbonium Ions with Nucleophiles’ Calvin D. Ritchie, G . A. Skinner, and V. G . Badding Contribution f r o m the Department of Chemistry, State University of New York at Buffalo, Buffalo, New York 14214. Received November 10, 1966
Abstract: The rates and equilibria of reactions of three derivatives of malachite green [4,4’-bis(dimethylamino)triphenylmethyltetrafluoroborate] with several nucleophiles in water, methanol, dimethyl sulfoxide, and dimethyl-
formamide solvents have been studied. In all solvents azide ion reacts much faster with the carbonium ions than does cyanide ion, although the equilibrium constants for formation of the covalent product are greater in the case of cyanide. The rate constants for reactions of the carbonium ions with the nucleophiles are found to be more sensitive to changes in solvent than are the equilibrium constants. Both forward and reverse rate constants are changed in the same direction by a change of solvent. It is argued that this latter observation requires the consideration of solvent reorientation as an essential part of the activation process.
-
deas concerning solvent effects on reactivity have undergone drastic revision since the classical concept of the solvent as a continuous dielectric medium. In spite of many recently proposed theories of the nature of the effect of solvents on reactivities in solution which focus attention on some particular aspect of the solvent such as ionizing power,2 dipolar associat i ~ n ,or~ hydrogen bonding a b i l i t ~ ,we ~ feel that a much more fundamental question requires attention. In somewhat naive terms, this question may be phrased as follows. Can the reactions of molecules in solution be discussed meaningfully in terms of only those molecules which are observed to undergo the chemical change, or is it necessary to take account of the entire system, including the solvent motion along the reaction coordinate? In studies on quite different systems, both Caldin5 and Robertson6 have been led to revive an earlier suggestion’ that solvent reorganization may make an appreciable contribution to energies of activation in solution. Essentially, the concept suggested is that initial-state solvation has been broken down to some extent without compensation by complete final-state solvation. This concept would require the consideration of the motion of solvent molecules in discussions of reactions in solution, and is counter to the common assumption that all solvent is in the most stable state at all points along the reaction coordinate. We have felt that pertinent information concerning the question posed above can best be gained by a study where the mechanism of reaction is least complicated. In the present paper, we report the results of a study of solvent effects on reactions of stabilized carbonium ions with the anionic nucleophiles cyanide, azide, hydroxide, and methoxide. Malachite green
[4,4’ bis(dimethy1amino)triphenylmethyl tetrafluoroborate] and its 4”-substituted derivatives, crystal violet [4,4’,4” tris(dimethy1amino)triphenylmethyl tetrafluoroborate)] and p-nitro malachite green [4,4’bis(dimethy1amino) - 4” - nitrotriphenylmethyl tetrafluoroborate], have been shown in a variety of studiess-14 to react with nucleophiles at the methyl carbon to form covalent derivatives. The reactions are generally kinetically straightforward and are easily followed by spectrophotometry at low concentrations because of the very high extinction coefficients ( E = lo5 M-’ cm-l) in the 590-650-mp range.
(1) This work was supported by a grant from Army Research Office (Durham), DA-ARO(D)-31-G376. (2) S. Winstein, A. H. Fainberg, and E. Grunwald, J . Am. Chem. Soc., 79, 4146 (1957); S. G. Smith, A. H. Fainberg, and S. Winstein, ibid., 83, 618 (1961). (3) R. W. Taft, Jr., G. B. Klingensmith, E. Price, and I. R. Fox, paper presented at Symposium on Linear Free Energy Correlations, Durham, N. C., Oct 1964; see also, J. Am. Chem. Soc., 87, 3620 (1965). (4) A. J. Parker, Quart. Reu. (London), 16, 163 (1962). (5) E. F. Caldin, J. Chem. Soc., 3345 (1959). (6) K. T. Leffek, R. E. Robertson, and S. Sugamori, J. Am. Chem. Soc., 87, 2097 (1965), and other references cited there. (7) E. A. Moelwyn-Hughes, Proc. Roy. SOC. (London), A164, 295 (1938); A212, 260 (1952).
(8) 0. F. Ginsberg and N. S. Melnkova, J . Gen. Chem. USSR, 25, 1109 (1955). (9) J. C. Turgeon and V. K. LaMer, J . Am. Chem. Soc., 74, 5988 (1952). (10) M. Gillois and P. Rumpf, Compt. Rend., 238, 591 (1954). (11) G. H. Brown, S. R. Adisesh, and J. E.Taylor, J . Phys. Chem., 66, 2426 (19621. ( 1 2 ) C . 6.Ritchie, W. F. Sager, and E. S. Lewis, J. Am. Chem. Soc., 84, 2349 (1962). (13) C. G. Ekstrom, Acta Chem. Scand., 20, 444 (1966), and earlier references cited there. (14) E. F. J. Duynstee and E. Grunwald, J . Am. Chem. Soc., 81, 4542 ( 1959). (15) R. W. Taft, Jr., personal communication.
I
-
Results The reactions of malachite green, p-nitro malachite green, and crystal violet with hydroxide ion and water in aqueous solution were studied in buffer solutions chloride under of triethylamine-triethylammonium pseudo-first-order conditions. In the case of crystal violet, the rates in the high pH range were studied in unbuffered solutions of standard potassium hydroxide. Plots of the observed pseudo-first-order rate constants us. concentrations of hydroxide ion are shown in Figures 1-3. From the slopes and intercepts of these plots, we obtain the values for the second-order rate constants for reaction with hydroxide ion, and the pseudofirst-order rate constants for reaction with water which are reported in Table I. The reactions with hydroxide ion and with water are found to be rather insensitive to inert salt effects. Addition of potassium perchlorate in concentrations M has no measurable effect on the rates of up to the reactions. Taft has similarly reported15 that the addition of potassium chloride in concentrations of 0.05-1.0 M causes less than a 10% change in the rate
~~
Ritchie, Skinner, Badding
Solvent Effects on Carbonium Ion-Nucleophile Reactions
2064
10.0
-'1
-
P *
*
8.0 6.0
4.0
Y
2.0
0.0
1 (OH-)
I
10'
Figure 1. Plot of pseudo-first-order rate constants for reaction of crystal violet with water and hydroxide ion in water: 0, present work; 0 , data from ref 15; A, data from ref 14.
10.0 I2.O h
'V
0
In
I
8.0
-
6.0
.
4.0
.
v
Q I
*
Y
2.0
i."" L
1.0
2.0
3.0
4.0
6.6
6.0
( O H - ) a IO'
Figure 2. Plot of pseudo-first-order rate constants for reaction of malachite green with water and hydroxide ion in water: 0, present work; 0 , data from ref 15; A, data from ref 14.
constant for the reaction of malachite green with hydroxide ion. The rate constants reported in Table I may be compared with a number of previously reported values. For the reaction of malachite green with hydroxide ion at 25", values of 1.5915 and 1.3614 M-' sec-' have Table I. Rate Constants for Reactions with Hydroxide Ion and Water in Aaueous Solution at 25.0" koa-, M-'
Cation Crystal violet Omax 590 m d Malachite green (hm.,615 m d p N i t r o malachite green Amax 641 mp) Crystal violet in D20 a
sec-1
kHpo X lo4,sec-1 pKo
2 . 0 5 X IO-'
0.35
9.36
2.18
2.11
6.84
5.64
1.80
5.50
2 . 2 x 10-1
0.35
...
Values taken from ref 8.
been reported. Values of the first-order rate constant for the reaction with water from the same sources are 2.2 X and 3.8 X sec-l, respectively. The Journal of the American Chemical Society j 89:9
April 26, 1967
1.0
2.0
3.0
4.0
5.0
6.0
( O H - ) x 10'
Figure 3. Plot of pseudo-first-order rate constants for reaction of pnitro malachite green with water and hydroxide ion in water.
discrepancies of these values with those reported in Table I are primarily due to the determination of the slope and intercept of the plot of observed rate constant us. concentration of hydroxide. The actually observed rates, as shown in Figures 1 and 2, are in reasonable agreement. The same explanation applies to the comparison of the rate constants reported for crystal violet. Ekstrom, l 3 however, has reported values considerably lower than these for all three of the carbonium ions studied here. His measurements were carried out at 20" rather than the 25" used in the present study, but it is most unlikely that a 5' temperature difference could explain the large discrepancies. The rates of reaction of the carbonium ions with cyanide ion in aqueous solution were determined by measuring the increase in rate of reaction of the dyes on increase of cyanide ion concentration at constant pH and buffer concentration. The pH of all solutions used in these determinations was accurately adjusted to 10.61. Since hydrogen cyanide has a pK of 9.14 in aqueous solution, only about 3 % of the cyanide ion added is converted to the conjugate acid. The rate constants obtained from these measurements were found to decrease with increasing cyanide ion concentration. The addition of potassium perchlorate to the solutions was also found to retard the reaction. The data are presented in Table 11. The salt effects noted are of the magnitude expected from Debye-Huckel theory for the neutralization of two unit charges. Rate constants extrapolated to zero ionic strength were calculated from a visual fit of the data plotted in Figure 4. The lines shown have slopes of -0.10 M-l. We have also attempted to measure the rates of reaction of the dyes with azide ion in aqueous solution, but have been prevented from doing so because of solubility problems. The equilibrium for formation of the triarylmethyl azide is too unfavorable to allow measurements at concentrations of less than loF2 M azide, and, at much higher concentrations of azide, salting out of the dyes is evident. The reactions of crystal violet with deuterioxide ion, deuterium oxide, and cyanide ion in DzO were studied by the same methods as outlined above for aqueous
2065 Table 11. Reactions of Carbonium Ions with Cyanide Ion in Aqueous Solution at 25.0" a
Cation Crystal violet
Malachite green
p-Nitro malachite green
Total kcN- X salt,b Cyanide, 10, M-1 M x 103 M x lo3 sec-1 8.33 8.33 8.33 10.01 10.8 13.3 15.0 16.6 18.3 20.8 23.3 8.33 10.3 18.3 23.3 3.66 3.83 4.33 5.33 6.00 6.66 8.33 10.0 11.7 13.3 15.0 8.33 13.3 4.33 5.33 6.66 8.33 10.0 11.7 13.3
px/zx
10.0 11.7 13.3 15.5 17.5 20.0 5.00 5.00 5.00 5.00 0.333 0.500 1.00 2.00 2.67 3.33 5.00 6.67 8.33 10.0 11.67 5.00
0.630 0,666 0.675 0.665 0.675 0.690 0.633 0.631 0.663 0.617 0.630 5.33 5.10 4.74 4.53 5.98 5.61 4.61 5.75 5.75 5.60 5.37 5.11 5.09 4.96 4.96 8.88
9.14 9.14 9.14 10.0 10.4 11.5 12.3 12.9 13.5 14.4 15.3 9.14 11.5 13.5 15.3 6.05 6.19 6.58 7.31 7.75 8.16 9.14 10.0 10.8 11.5 12.3 9.14
5.00 1.00 2.00 3.33 5.00 6.67 8.33 10.0
8.15 9.9 9.25 8.96 9.17 8.67 8.36 8.21
11.5 6.58 7.31 8.16 9.14 10.0 10.8 11.5
5.0 5.0 5.0 6.67 7.50
a Extrapolated rate constants: crystal violet, 8.8 X lo-? M-' sec-1; crystal violet in D20, 8.5 X M-' sec-I; malachite green, 6.9 X 10-l M-' sec-'; p-nitro malachite green, 1.13 M-' sec-1. M buffer salt, plus Total salt concentration is 3.33 X potassium cyanide concentration shown, plus potassium perchlorate.
solution. The values obtained for the rate constants are reported in Tables I and 11. The solvent isotope effect on all three of these reactions is remarkably small. All of the reactions discussed above were conveniently studied by conventional spectrophotometric techniques, the half-times for the reactions varying from several minutes to several hours. In all of the studies in nonaqueous solutions discussed below, however, the reactions are too fast for conventional techniques, and we have used a stop-flow apparatus patterned after the one described by Gibson.lG In our apparatus, the solutions contact only glass, Teflon, and stainless steel. The apparatus is described in the Experimental Section. Before presenting the results obtained in the nonaqueous solvents, we would like to emphasize the extreme importance of rigorously purified Solvents, and Of carrying Out Operations in an inert, dry atmosphere. Contamination of solutions by moisture or traces of acids can cause astonishing decreases in the rates of some of these reactions, particularly those with cyanide ion. (16)
Q.H.Gibson and L. Milnes, Biochem. J., 91,
1.00
102, M
161 (1964).
I
a
'
0.60 6.0
8.0
12.0
10.0
14.0
16.0
18.0
Figure 4. Plot of log kc"- us. square root of ionic strength for reaction of carbonium ions with cyanide ion in water: 0, malachite green (1 log k is shown); A, crystal violet (2 log k is shown); 0 , p-nitro malachite green (1 log k is shown).
+
+
+
The reactions of the three carbonium ions with methoxide ion in anhydrous methanol were studied by reaction of solutions of the dye with dilute solutions of standardized sodium methoxide. The rates of the reactions showed a very strong salt effect at sodium M . It was methoxide concentrations above found that the assumption of a dissociation constant M brings for sodium methoxide ion pair of 1.03 X the rate constants obtained over a wide range of concentrations into agreement with each other. Since at least one determination of rate was made at concentrations of sodium methoxide below M for each dye, the accuracy of the reported rate constants is not very dependent on this choice for the ion pair dissociation constant. The values obtained for the rate constants are shown in Table 111. Table 111. Rate and Equilibrium Constants in Methanol Solution at 25.0' a Cation Malachitegreen Crystal violet p-Nitromalachite green m-Chloro malachite green m-Nitro malachite green
kocH3-
X 10-8 PKR'
kx3-d
KNa-e
1.25 8.80 ... 0.065 . . . ... >10-2 6.0 6.79 6 . 2 X lo4 6 . 2 X 2.9
7.75
7.4
6.84
.. ...
...
a pKBx3 in methanol = 8.91; pK (triethylammonium ion) in methanol = 10.88; pK (tribenzylammonium ion) in methanol = 6.40; A,, in methanol: crystal violet, 592 mp; malachite green, 623 mp; p-nitro malachite green, 648 mp. Rate constant in M-1 sec-1 for the reaction with methoxide ion. Equilibrium constant for the reaction Ar3C+ MeOH ArlCOMe H+ at 25.0". d Rate constant in M-1 sec-1 for reaction with azide ion. e Equilibrium constant for the reaction Ar3CN3 Ar3C+ N3-.
+
+ +
In a previous stUdy,~2we have reported the rates of reactions of several malachite green derivatives with methoxide ion using rates relative to the basicity of triethylamine in methanol. F~~~ the present data in conjunction with the previous measurements, therefore, we can evaluate the pKb of triethylamine as 6.04. Using the solvent ionization constant of pK, = 16.92
Ritchie, Skinner, Badding
1 Solvent
Effects on Carbonium Ion-Nucleophile Reactions
2066
reported by Koskikallio, li we thus obtain a pK, value of 10.88 for triethylammonium ion in methanol. In water, the pKa is 10.87,'* and, as we have pointed out previou~ly,'~ would be expected to remain the same in methanol. In the earlier study, we found that the rate of reaction of the carbonium ions with methanol is negligible, and this conclusion is confirmed by the present study. Also in the earlier work,'Z we determined the PKR values of several malachite green derivatives relative to tribenzylammonium ion in methanol. We have now measured the pKa of tribenzylammonium ion in methanol by the differential potentiometric method described in another paperIg and find a value of 6.40 at 25.0". This value may now be used to convert the previously reported relative pKR values to absolute values. These values are shown in Table 111. The reaction of azide ion with p-nitro malachite green in methanol was studied by observation of the first 50-100 msec of reaction of the carbonium ion M sodium with a methanol solution of 3.12 X M methoxide with 1.04 X 10-2 to 9.73 X added sodium azide. The reaction with azide under these conditions is fast enough that equilibrium is established before any appreciable reaction with methoxide ion occurs. From the stop flow traces shown in the Experimental Section, both the equilibrium constant and rate constant for the formation of the triarylmethyl azide can be evaluated. The values obtained for these constants are shown in Table 111. The equilibrium constants for formation of the triarylmethyl azides from malachite green and crystal violet are too small to allow measurements even at concentrations of lo-* M in sodium azide. We have measured the ionization constant for hydrazoic acid in anhydrous methanol by titration of sodium azide with methanolic perchloric acid using the potentiometric method described elsewhere. 2o We obtain a pK value of 8.9 at 25.0°, which is in excellent agreement with the value of 8.9 reported by Parker.21 In this work, in order to provide a convenient buffer solution for calibration of glass electrodes in methanol, M we have measured a pH of 8.987 for a 1.465 X solution of potassium hydrogen phthalate in anhydrous methanol at 25.0". Unfortunately, we have not been able to measure the rates of reactions of the carbonium ions with cyanide ion in methanol because of the basicity of cyanide ion in this solvent. From titrations of tetraethylammonium cyanide in methanol, we have estimated that Kb for cyanide is approximately M . We cannot, therefore, prepare solutions of reasonable concentration in which the amount of cyanide exceeds the amount of methoxide present. On the basis of one experiment, however, it appears that cyanide reacts more slowly than does methoxide. A solution conM sodium methoxide and 5.5 X taining 7.8 x 10-3 M tetraethylammonium cyanide was found to react with crystal violet at the same rate as does the
The reactions of all three carbonium ions with azide ion in dimethylformamide solution were also found to reach equilibrium on mixing. Consistent equilibrium constants over a range of concentrations were again obtained, and lower limits on the rates of reactions set at lo7 M-I sec-'. The reactions of the dyes with cyanide ion in dimethylformamide were studied in the same manner as
(17) J. Iio7
6.1 x 10-4 5.0 X 10-6
8.91 7.9
1.13
DMF > MeOH > HzO, which is the expected order for the dispersion effect. It seems probable that this effect is
/ April 26, 1967
(27) E. M. Arnett, Progr. Phys. Org. Chem., 1, 223 (1963). (28) E. Grunwald and E. Price, J . A m . Chem. Soc., 86, 4517 (1964)
2069
primarily responsible for the differences in rates and equilibria between dimethyl sulfoxide and dimethylformamide solvents, since electrostatic effects are expected to be quite similar in the two solvents. The order of nucleophilicities of hydroxide, azide, and cyanide ion toward the carbonium ions studied here is the reverse of the usual nucleophilicities in s N 2 type reactions. In the present work, azide reacts faster than does methoxide, which is faster than cyanide. In ~ cyanide reacts about eight times typical S N reactions, faster than hydroxide, and hydroxide reacts about twice as fast as does azide.29 Apparently, polarizability and the a effect are primary factors in nucleophilicity toward these stabilized carbonium ions. It is also interesting to note that the relative nucleophilicities of azide and cyanide ions are not changed by solvent. Azide ion appears to react at least lo4 times more rapidly than does cyanide in all of the solvents studied. It also appears, however, that the relative acidities of hydrazoic acid and hydrocyanic acids are not drastically changed by a change of solvent. In water, the difference in pK's is 4.4 units, and in dimethyl sulfoxide, the difference is 5.0 units. One would suspect, therefore, that there is no great difference in the hydrogen bonding of water to azide and cyanide ions. It is apparent, then, that the greater reactivity of azide ion is not due to solvent hydrogen bonding strengths. Experimental Section Materials. The tetrafluoroborate salts of malachite green, p-nitro malachite green, and crystal violet were prepared as previously described.12 After extensive recrystallization from water, the compounds exhibited identical spectral properties with those previously reported by us, and in close agreement with those reported by ot1iers.S-14 Elemental analysis of malachite green tetrafluoroborate gave 66.59% C, 6.10% H, and 6.69% N (theoretical for C?3H2aN2BF4:66.4% C, 6.01 % H, and 6.73 N). Triethylamine (Eastman White Label) was used without further purification. Tri benzylamine was recrystallized from ethanol before use. Sodium azide (Eastman Practical Grade) was recrystallized several times from ethanol and dried at 120" for several hours. Potentiometric titration of the salt with perchloric acid in methanol solution required 98.5% of the theoretical amount of acid. Potassium cyanide (Baker Analyzed Reagent) was used without further purification. Tetraethylammonium cyanide30 was recrystallized from dry acetonitrile under an argon atmosphere and dried under vacuum at room temperature. Potentiometric titration of the salt with p toluenesulfonic acid in dimethyl sulfoxide solution required precisely the thcoretical amount of titrant. The dried salt was stored under argon in a desiccator. Sadium perchlorate and potassium perchlorate were obtained in reagent grade quality and were used without further purification. Hydrochloric acid solutions were standardized by potentiometric titration of analytical grade sodium carbonate in water. Potassium hydroxide solutions were standardized by potentiometric titration of a National Bureau of Standards sample of potassium hydrogen phthalate. Sodium methoxide solutions were prepared by reacting freshly cut sodium with methanol purified as described below. The solutions were standardized by potentiometric titration with perchloric acid in methanol which, in turn, had been standardized with potassium hydrogen phthalate in methanol.19 All operations were carried out under an inert atmosphere, and the solutions were stored under argon. Solvent Purification. All solvents were purified and stored under a dry inert atmosphere. Water was doubly distilled from potassium permanganate.
Methanol was purified by the ion-exchange technique previously described. 19 Potentiometric titration of the solvent with methoxide and with perchloric acid showed much less than 5 X 10-6 M acidic or basic impurities. Karl Fisher titration shows less than 10 ppm water present. Deuterium oxide (Columbia Organic Chemicals Co., 99.5 % isotopic purity) was purified by an ion-exchange technique similar to that described for methanol. The solvent used was shown to contain less than 5 x 10-6 M acidic or basic impurities, and nearinfrared analysis showed better than 99.5 % isotopic purity. Dimethylformamide was dried with molecular sieves and vacuum distilled from phosphorus pentoxide as previously described. 2 4 Potentiometric titration of the solvent with methanolic sodium methoxide and with picric acid in dimethylformamide reveals less than 5 x 10-6 M acidic or basic impurities. Karl Fisher titration shows less than 10 ppm water present. The solvent was used within 48 hr following preparation. Dimethyl sulfoxide was purified by a procedure furnished by Steiner,sl in which the solvent is dried with molecular sieves, topped by distillation under vacuum, and finally distilled from freshly prepared potassium amide under vacuum in a rotary evaporator. The distillation temperature is kept well below 50". The distillate is held under vacuum for several hours t o remove traces of ammonia. Potentiometric titration of the solvent with dimsyl ion solutions and with p-toluenesulfonic acid in diM acidic or basic methyl sulfoxide22 shows less than 5 X impurities, and Karl Fisher titration shows less than 10 ppm water. Apparatus. A Cary Model 14 spectrophotometer equipped with a specially designed thermostated cell holder was used to record spectral data and for kinetic measurements on those reactions with half-lives greater than about 1 min. The thermostated cell holder consisted of a brass block through which thermostated water was circulated. Cylindrical spectrophotometer cells of 0.01-10.0 cm fit snugly into a bore of the block. Checks of the temperature of solutions in the cells showed that a temperature variation of less than 0.1 from that of the circulating liquid could be maintained between 5 and 35". A Beckman KF3 aquameter was used for the Karl Fisher titrations. Potentiometric titrations and pH measurements were carried out by the use of either a Beckman 1019 research pH meter for nonaqueous solutions, or a Beckman Model G p H meter for aqueous solutions. Measurements in aqueous solutions utilized a Beckman Type E2 glass electrode and a saturated calomel reference electrode. The electrode systems for use in nonaqueous solutions have been described elsewhere. 19,22.21 For the study of the kinetics of reactions having half-lives of less than 1 min, a small-scale stop-flow apparatus was used. The apparatus was patterned closely after the one described by Gibson and Milnes. l6 Hamilton Co. gastight syringes (2.5 ml) were used for the driving syringes. Seals between various parts of the apparatus were made with Teflon gaskets kept under pressure. Valve tips and mixing chambers were also fabricated from Teflon. In the finished apparatus, the reaction solutions contact only glass, Teflon, and stainless steel. (It is important that no epoxy seals are used since they are rapidly softened by the solvents used in the present study. Lucite is, of course, also unsuitable for use with these solvents. We have had some indications that Kel-F is attacked to some extent by strongly basic solutions in dimethyl sulfoxide.) The observation chamber of the apparatus has a 2.0-cm light path 2.0 mm in diameter fitted with Teflon-gasketed quartz windows. This chamber is interchangeable with one with a 2.0-mm light path. The light source is a Bausch and Lomb high-intensity grating monochromator, Model No. 33-86-25-02, with beam condenser, modified to use a G E No. 1958 tungsten-iodine high-intensity lamp, and to operate from a stabilized D C power supply (Electro, Model NFB). A Corning filter, (33-74, was used to remove higher order spectra. Light intensity was measured with a 1P-28 phototube whose output was fed through a cathode follower-emitter (R. W. Dykstra, Sauk Village, Ill.) to an oscilloscope (Tektronix Type 564 MOD08, time base No. 2B67, differential amplified No, 2A63). Each kinetic run on the stop-flow apparatus requires approximately 0.5-1.0 ml of solution after the apparatus has been thoroughly flushed with the reactants. A total of about 10.0 ml of each reactant solution is generally sufficient for flushing and
(29) J. 0. Edwards and R. G . Pearson, J . Am. Chem. Soc., 84, 16 ( 1962).
(30) We are grateful to Drs. S.Andreades and 0. Webster of DuPont for a generous sample of this material and for purification procedures.
(31) E. C. Steiner, to be published. We are grateful to Dr. Steiner for permission to quote this procedure.
Ritchie, Skinner, Badding / Solvent Effects on Carbonium Ion-Nucleophile Reactions
2070 several measurements. Mixing time in the apparatus is estimated t o be less than 5 msec with a good sharp push on the driving block. The entire valve block, delivery block, driving syringes, and observation chamber are thermostated by circulating water from a constant temperature bath. Kinetics in Aqueous Solutions. The rates of all the reactions of malachite green and p n i t r o malachite green, and some of the reactions of crystal violet, with water and hydroxide ion were studied in solutions buffered with triethylamine and its hydrochloride salt. The usual technique was to prepare a solution ca. 0.25 M in amine and a separate solution of 0.25 M hydrochloric acid. These solutions were then mixed, with dilution, to produce a solution of the desired pH and a total buffer concentration of ca. 10-2 M (amine amine hydrochloride). The solution was then placed in a thermostat set at 25.0". A measured volume of this solution was then mixed with an equal volume of a thermostated solution of the carbonium ion. One portion of the resulting solution was placed in the spectrophotometer cell and the absorbance of the solution at the wavelength of maximum absorbance of the dye followed with time. The p H of the remaining portion was measured at the start and at the end of the kinetic run. All solutions were kept under a nitrogen atmosphere. The initial dye concentrations ranged from 3 x 10-6 to 5 x 10-6 M. In the case of crystal violet, a number of runs were made in the high pH region by use of standard solutions of potassium hydroxide. The p H of the solutions was measured as described above. The second-order rate constants for reactions with hydroxide ion, and the pseudo-first-order rate constants for reaction with water were determined from the slopes and intercepts, respectively, of plots of the observed pseudo-first-order rate constants us. concentration of hydroxide ion. Figures 1, 2, and 3 show the plots obtained. The lines shown in the figures were determined by the method of least squares. A fivefold increase in buffer concentration was found to cause about a 10% increase in the observed rate at a given pH. We believe that this is due to the presence of small amounts of impurities in the triethylamine. The effect of inert salts on the rates of the reactions was examined by adding up to M potassium perchlorate t o solutions at constant p H and buffer concentration. No measurable effect was observed for any of the reactions. The reactions of cyanide ion with the dyes were studied by adding measured quantities of potassium cyanide t o solutions containing a constant buffer concentration and a constant p H of 10.61. The rate constant for reaction with cyanide was obtained by subtraction of the pseudo-first-order rate constant observed in the absence of cyanide from that in the presence of cyanide. The effect of inert salts was examined by adding various amounts of potassium perchloiate t o solutions of constant pH, buffer, and cyanide ion concentrations. Typical results of these experiments are shown in Table I1 and Figure 4. In the aqueous solutions, precipitation of the dyes at salt concentrations much above M prevented further examination of salt effects. The same factor prevented the study of the reactions of the dyes with azide ion in aqueous solutions. At concentrations M , it appeared that there was a slight reaction of azide about with p-nitro malachite green, but precipitation of the dye was enough of a problem that accurate measurements could not be made. A very rough estimate for the dissociation constant of p-nitro malachite green azide in water can be given as 5 X loW2M . We would not be surprised, however, if this value is as much as a factor of 10 too low; it is unlikely that it is too high. Reactions of Crystal Violet in D20. The reactions of crystal violet with deuterium oxide and deuterioxide ion in deuterium oxide were studied in unbuffered solutions of sodium deuterioxide. Deuterioxide ion concentrations ranged from 2 X to 8.5 X 10-3 M . Treatment of the data was as described for aqueous solutions. The rate constant for reaction with cyanide ion in D20 was determined in the same manner as for the reaction in water. The concentration of deuterioxide ion in these runs was 1.61 X lo-* M. Reactions in Methanol Solution. The reactions of the dyes with methoxide ion in methanol were studied in unbuffered solutions of and sodium methoxide at concentrations between 8 X 7.8 X 10-8 M. The reactions showed half-times of ca. 0.1 t o 10 sec, and were studied on the stop-flow apparatus. Solutions were prepared and transferred t o gastight syringes in an argon-filled glove bag. The gastight syringes fit into adapters on the valve block of the stop-flow apparatus, and transfer of solutions could be accomplished without opportunity for contamination from the
+
Journal of the American Chemical Society
89:9
atmosphere. The stop-flow apparatus is thoroughly flushed with the reactants by twice filling the driving syringes and emptying through the flow path. The time base of the oscilloscope was adjusted during the flushing operations so that approximately 95 of the reaction was followed on a single trace. The second-order rate constants obtained for the reactions with methoxide ion were found t o decrease with increasing methoxide concentration. Addition of sodium perchlorate to the reaction solutions was also found to retard the rates. The assumption iM for sodium of an ion-pair dissociation constant of 1.03 X methoxide was found to be consistent with the data. A typical set of data for the reaction of malachite green with methoxide is shown in Table VII. Table VII. Reaction of Malachite Green with Methoxide Ion in Methanol at 25.0" Na,+ M X
lO3a
1.56 0.624 0.312 0.156 2.67 5.43 7.80
OCHa-, M X lo3* 1.56 0.624 0.312 0.156 1.56 1.56 7.80
kobsdrC
10-3
sec-1
k2("bsd)d
1.79 0.686 0.378 0.180 1.52 1.18 6.61
1.15 1.10 1.21 1.15 0.974 0.757 0.848
10-3 k2(cor)e 1.31 1.17 1.26 1.19 1.20 1.12 1.27
Total sodium ion concentration added. * Total sodium methoxide concentration. Pseudo-first-order rate constant obtained from plot of log A us. time. d Second-order rate constant obtained by dividing kl by total sodium methoxide concentration. Units are M-1 sec-'. e Second-order rate constant obtained by dividing kl by the concentration of free methoxide ion calculated from a dissociation constant for the sodium methoxide ion pair of 1.03 X 10-2 M. Units are M-l sec-1. The reaction of crystal violet with a solution of 7.8 X 10-3 M sodium methoxide and 5.5 X 10-3 M tetraethylammonium cyanide was found to proceed at the same rate as that in the absence of the cyanide. The reaction ofp-nitro malachite green with azide ion in methanol was studied in solutions containing 3.12 X IO-' M sodium methoxide. The reaction with methoxide at this concentration has a half-time of 0.37 sec. It was observed that when sodium azide is added to the solution, the initial absorbance shown on a time scale of 0.2 sec/div on the oscilloscope was decreased from that observed in the absence of azide, and that further decrease in absorbance is slower than in the absence of azide. The changes in initial absorbance with changes in azide concentration observed M in these traces gave an equilibrium constant of 6.2 (=k0.3)X for the dissociation constant of the triarylmethyl azide. In the absence of azide ion, a decrease in absorbance from 1.310 to 1.191 is observed in the first 50 msec of reaction ofp-nitro malachite green with the 3.12 X M methoxide solution. It was, therefore, possible to observe the reaction of azide with the dye, without appreciable interference from the simultaneous reaction with methoxide, by the following general technique. The vertical scale of the oscilloscope was adjusted so that 0.0-0.20 v corresponded to O-loO% transmittance. A trace of the reaction was then obtained on the usual time scale (usually 0.2--0.5 sec/div) and the initial transmittance noted. The vertical oscilloscope base was then adjusted to 20 mv for full-scale deflection, and the time base was adjusted to 5-10 msec/div. The base-line output was adjusted by means of a vernier control incorporated into the emitter follower circuit so that the zero line on the oscilloscope corresponded to the previously observed initial transmittance. A second trace of the same reactants was then obtained. From this latter trace, the second-order rate constant for the reaction of azide could be obtained by the usual method for pseudo-first-order reversible reactions. A typical example OF the traces obtained is shown in Figures 5 and 6 . The two traces in Figure 6 are for consecutive traces on the same reactants. The flat portion of the fast traces represents the portion of continuous flow immediately preceding the stop. In attempts to measure the reaction of malachite green with sodium azide in methanol, it was found that sodium azide concentrations up to 1.26 X M caused no decrease in the initial absorbance of the dye, and that the rate of reaction with methoxide
April 26, I967
Figure .5. Oscilloscope trace for slow time scale o f reaction of p nitro malachite green with methoxide ion in methanol in the presence of azide ion.
nas affected only to the extent calculated lor the eKect of the added sodium ion on the formation of sodium methoxide )on pair. Reactions of cr)staI vinlrt with azide ion i n methanol were not attempted. R e a d i m of Cyanide and Adde in Dinethylfmamide and Dimethyl Sulfoxide. The general technique for the stud) of reactions in dimethylformamide and in dimeth)l sulfoxide was as follows. Ma\ler solutions o f the d)e (m.10-6 M ) and of the salt (sodium azide or tetraeihylammonium cyanide) uere prepared h) weighing the sample on a Cahn electrobalancc and dissolving in a measured volume of the solrent. The reactions were studied in a manner analogous to that described above for methanol solutions. All reactions u e r e performed under pseudo-first-orderconditions. and tenfold range, of concentrations of nucleuphiler were studied to verif) the order of the reactions. For the a i d e reactions in both dimethylformamide and dimethyl sulfoxide, the solutions had reached equilibrium before stopping, even in the lonest concentrations of azide which would cause mrahurable changes In absurbdncc. Consistent values for the equilibrium constants were obtained over a range o f acidc concentrations in ever) case. Minimum values for the forward rate constants were calculated on the basis of mixing requiring 5 msec, and that the reactions uere 75": complete in this time. Actually, absolutely no change in absorbance could be ohsewrd in any case, and the reaction uas quite probably more than 95% complete uithin 5 mscc. Thus, the lower limils reporled m Tables I V and V arc \cry conservative values. In dimethyl sulfoxide solulion, reaction o f crystal violet with IO-> M azide solutions could not be detected. The limited solubility of sodrum azide m pure dimethyl sulfuxide" pretented studies at higher concentrations. The reactions of the dyes u i t h cyanide ion in both solvents were found to be complete at the concentrations of IO-' to IO-' M cyanide n hich were studied. We have found that the slightest contamination of dimethyl sulfoxide or dimethylformamide solutions 01c)anide by exposure to the atmosphere rcsull%in drastically reduced rates of reactions with thed)es.
Figure 6. Oscilloscope trace for fast time scale of reaction o f p nitro malachite green with azide ion in methanol. Since &rnerhyl sulfoxide is known 10 carry oarious marerials in solurion rhrough rhe skin, exrreme caurion should be exercised in handlingsolutions of azide or cyanide in rhis solvent. We ham used rubber gloues in handling conroiners wirh rhese solurions, ami h o e raken specialpains 10 clean up rhoroughly afrer each experimenr. pKMeasurements in NonaqueousSolvents. The measurementsof pK values for various acids reported in the results section have been camed out hy methods described in other papers from our laboratories."-m. 2 2 . 1 4
Acknowledgments. Although primary financial support of this work was through the Army Research Office (Durham), as acknowledged in footnote 1, we also wish to express our gratitude to the National Science Foundation (GP-2635) and the Public Health Service (GM12832) for grants supporting other phases of our work which resulted in many of the solvent purification techniques, pK measurement methods, and in construction of the stop-flow apparatus, which were necessary parts of the present research project. In addition, thanks are due to Mr. Htain Win for preparation of the carbonium ion salts and some early exploratory measurements in nonaqueous solutions; to Mr. G. H. Megerle for help in the purification of dimethylformamide; and to Mr. R. E. Uschold for help in the purification of dimethyl sulfoxide. We are also grateful to Professor R. G. Wilkins for helpful discussions on the construction of the stop-flow apparatus. The Cary Model 14 spectrophotometer used in the present work was purchased through Grant No. GM11036 from the Public Health Services. The Graduate School, State University of New York at Buffalo, furnished funds for the purchase of the Beckman KF3 aquameter.
Ritchie, Skinner, Badding 1 Solvent Effects on Carbonium IorrNucleophile Rewlions
